The hydrationequilibrium constant at 25 °C is called Kh, which in the case of carbonic acid is [H2CO3]/[CO2] ≈ 1.7×10−3 in pure water[3] and ≈ 1.2×10−3 in seawater.[4] Hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO2 molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2 + H2O → H2CO3) and 23 s−1 for the reverse reaction (H2CO3 → CO2 + H2O). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. The addition of two molecules of water to CO2 would give orthocarbonic acid, C(OH)4, which exists only in minute amounts in aqueous solution.

Carbonic acid is an intermediate step in the transport of CO2 out of the body via respiratory gas exchange. The hydration reaction of CO2 is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which both increases the reaction rate and dissociates a hydrogen ion (H+) from the resulting carbonic acid, leaving bicarbonate (HCO3−) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO2 and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.[5]

The oceans of the world have absorbed almost half of the CO2 emitted by humans from the burning of fossil fuels.[6] The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels.[7] This process is known as ocean acidification.[8]

Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution, carbonic acid exists in equilibrium with carbon dioxide, and the concentration of H2CO3 is much lower than the concentration of CO2. In many analyses, H2CO3 includes dissolved CO2 (referred to as CO2(aq)), H2CO3* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows:[2]

Whereas this apparent pKa is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO2-containing solutions. A similar situation applies to sulfurous acid (H2SO3), which exists in equilibrium with substantial amounts of unhydrated sulfur dioxide.

At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2 solution) is completely determined by the partial pressure of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2CO3, HCO3− and CO32−) as well as of the hydration equilibrium between dissolved CO2 and H2CO3 with constant (see above) and of the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution:

The corresponding equilibrium equations together with the relation and the charge neutrality condition result in six equations for the six unknowns [CO2], [H2CO3], [H+], [OH−], [HCO3−] and [CO32−], showing that the composition of the solution is fully determined by . The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

(atm)

pH

[CO2](mol/L)

[H2CO3](mol/L)

[HCO3−](mol/L)

[CO32−](mol/L)

1.0 × 10−8

7.00

3.36 × 10−10

5.71 × 10−13

1.42 × 10−09

7.90 × 10−13

1.0 × 10−7

6.94

3.36 × 10−09

5.71 × 10−12

5.90 × 10−09

1.90 × 10−12

1.0 × 10−6

6.81

3.36 × 10−08

5.71 × 10−11

9.16 × 10−08

3.30 × 10−11

1.0 × 10−5

6.42

3.36 × 10−07

5.71 × 10−10

3.78 × 10−07

4.53 × 10−11

1.0 × 10−4

5.92

3.36 × 10−06

5.71 × 10−09

1.19 × 10−06

5.57 × 10−11

3.5 × 10−4

5.65

1.18 × 10−05

2.00 × 10−08

2.23 × 10−06

5.60 × 10−11

1.0 × 10−3

5.42

3.36 × 10−05

5.71 × 10−08

3.78 × 10−06

5.61 × 10−11

1.0 × 10−2

4.92

3.36 × 10−04

5.71 × 10−07

1.19 × 10−05

5.61 × 10−11

1.0 × 10−1

4.42

3.36 × 10−03

5.71 × 10−06

3.78 × 10−05

5.61 × 10−11

1.0 × 10+0

3.92

3.36 × 10−02

5.71 × 10−05

1.20 × 10−04

5.61 × 10−11

2.5 × 10+0

3.72

8.40 × 10−02

1.43 × 10−04

1.89 × 10−04

5.61 × 10−11

1.0 × 10+1

3.42

3.36 × 10−01

5.71 × 10−04

3.78 × 10−04

5.61 × 10−11

We see that in the total range of pressure, the pH is always largely lower than pKa2 so that the CO32− concentration is always negligible with respect to HCO3− concentration. In fact CO32− plays no quantitative role in the present calculation (see remark below).

For vanishing , the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO3− form.

For normal atmospheric conditions ( atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO2 form. From this pressure on, [OH−] becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H+ and HCO3−.

For a CO2 pressure typical of the one in soda drink bottles ( ~ 2.5 atm), we get a relatively acid medium (pH = 3.7) with a high concentration of dissolved CO2. These features contribute to the sour and sparkling taste of these drinks.

Between 2.5 and 10 atm, the pH crosses the pKa1 value (3.60) giving a dominant H2CO3 concentration (with respect to HCO3−) at high pressures.

A plot of the equilibrium concentrations of these different forms of dissolved inorganic carbon (and which species is dominant), as a function of the pH of the solution, is known as a Bjerrum plot.

Remark

As noted above, [CO32−] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H+]:

Theoretical calculations show that the presence of even a single molecule of water causes carbonic acid to revert to carbon dioxide and water. In the absence of water, the dissociation of gaseous carbonic acid is predicted to be very slow, with a half-life of 180,000 years.[9] This may only apply to isolated carbonic acid molecules however, as it has been predicted to catalyze its own decomposition[10]

It has long been recognized that pure carbonic acid cannot be obtained at room temperatures (about 20 °C or about 70 °F). It can be generated by exposing a frozen mixture of water and carbon dioxide to high-energy radiation, and then warming to remove the excess water. The carbonic acid that remained was characterized by infrared spectroscopy. The fact that the carbonic acid was prepared by irradiating a solid H2O + CO2 mixture may suggest that H2CO3 might be found in outer space, where frozen ices of H2O and CO2 are common, as are cosmic rays and ultraviolet light, to help them react.[9] The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuo, which causes protonation of bicarbonate, followed by removal of the solvent. The previously suggested alpha-carbonic acid, which was prepared by the same technique using methanol rather than water as a solvent was shown to be a monomethyl ester CH3OCOOH.[11]